A buffer is a solution that holds its pH steady when small amounts of acid or base are added. It contains a weak acid and its conjugate base in comparable amounts: the base form mops up added acid and the acid form neutralizes added base. Blood, seawater and almost every biological experiment depend on buffers. This calculator finds a buffer’s pH from its composition, shows how far the pH moves when you add strong acid or base, and works in reverse to give you a recipe, including the grams of each salt to weigh, for any target pH.
How to use the buffer pH calculator
- Choose what you need: the pH of a buffer, or a Recipe for a target pH.
- Select a Buffer system — formate, lactate, acetate, MES, carbonic acid/bicarbonate, phosphate, HEPES, Tris, ammonium/ammonia or bicarbonate/carbonate — or choose “Custom pKa”.
- For the pH, enter the Weak acid concentration [HA] and the Conjugate base concentration [A⁻] in the finished buffer. Optionally choose to add a strong acid or base, with the amount and the buffer volume.
- For a recipe, enter the Target pH, the Total buffer concentration and the Buffer volume.
- Read the pH (or the base-to-acid ratio), the buffer capacity and effective range, and for recipes the concentrations, millimoles and grams of each form.
Henderson–Hasselbalch equation
Adding x moles of strong acid converts x moles of A⁻ into HA; adding strong base does the reverse. Because only the ratio of the two amounts matters, diluting a buffer barely changes its pH. Buffer capacity is β = 2.303([H⁺] + [OH⁻] + CKa[H⁺] ÷ (Ka + [H⁺])²).
Worked example
An acetate buffer under attack
A buffer contains 0.10 M acetic acid and 0.15 M sodium acetate (pKa 4.76).
pH = 4.76 + log10(0.15 ÷ 0.10) = 4.76 + 0.176 = 4.94
Add 10 mmol of HCl to 1 L: acetate falls from 150 to 140 mmol and acetic acid rises from 100 to 110 mmol, so pH = 4.76 + log10(140 ÷ 110) = 4.86, a drop of only 0.07. The same acid in 1 L of pure water would give pH 2.00.
Designing a buffer. For 1 L of 0.10 M acetate buffer at pH 5.00, the ratio is 100.24 = 1.738. That splits into 0.0365 M acetic acid and 0.0635 M acetate: weigh 2.193 g of acetic acid (about 2.09 mL of glacial acid) and 5.207 g of anhydrous sodium acetate. A 0.5 L phosphate buffer at pH 7.40 and 0.10 M needs 2.321 g of NaH2PO4 and 4.352 g of Na2HPO4.
Common buffer systems
| Buffer | pKa (25 °C) | Useful pH range |
|---|---|---|
| Formate | 3.75 | 2.8–4.8 |
| Acetate | 4.76 | 3.8–5.8 |
| MES | 6.10 | 5.5–6.7 |
| Carbonic acid/bicarbonate | 6.35 | 5.4–7.4 |
| Phosphate | 7.20 | 6.2–8.2 |
| HEPES | 7.50 | 6.8–8.2 |
| Tris | 8.07 | 7.1–9.1 |
| Ammonium | 9.25 | 8.3–10.3 |
| Carbonate | 10.33 | 9.3–11.3 |
The masses assume anhydrous salts. If your bottle is a hydrate, such as sodium acetate trihydrate or disodium phosphate heptahydrate, scale the mass by the ratio of the molar masses.
Making a buffer in practice
Dissolve both components in about 80% of the final volume, measure with a calibrated pH meter, adjust with small additions of strong acid or base, and only then dilute to the mark. Calculated recipes get you close, but ionic strength and temperature shift the real pH slightly. Blood relies on the bicarbonate system. At body temperature its effective pKa is about 6.1 rather than the 25 °C value of 6.35, and a normal ratio of about 20:1 bicarbonate to dissolved CO₂ puts the pH near 7.4.
The pH calculator handles plain acid or base solutions, the titration calculator shows the buffer region on a full curve, and the molarity calculator converts the concentrations into masses for other compounds.
Frequently asked questions
What is the Henderson–Hasselbalch equation?
pH = pKa + log₁₀([A⁻] ÷ [HA]), where [HA] is the weak acid and [A⁻] its conjugate base. When the two are equal, the log term is zero and the pH equals the pKa. Each tenfold change in the ratio shifts the pH by one unit.
How do I choose a buffer system?
Pick one whose pKa is within one unit of the pH you want, ideally as close as possible. Acetate (4.76) covers about pH 3.8–5.8, phosphate (7.20) about 6.2–8.2 and Tris (8.07) about 7.1–9.1. Outside pKa ± 1 the buffer has little capacity on one side.
What is buffer capacity?
How much strong acid or base a buffer can absorb per unit change in pH. It is highest when [HA] = [A⁻] and grows with the total concentration. A 0.25 M acetate buffer at pH 4.94 has β ≈ 0.14 mol/L per pH unit, so about 14 mmol of NaOH would raise 1 L by 0.1.
How accurate is the Henderson–Hasselbalch equation?
It is very good between roughly pH 3 and 11 when both forms are above about 1 mM. In dilute solutions or at extreme pH, the acid's own dissociation and water matter; the calculator then shows an exact pH from the full proton balance. At high ionic strength, activities make real pH differ by a few hundredths to tenths.
Does temperature affect buffer pH?
Yes, because pKa changes with temperature. Tris is the best-known example: its pKa drops by about 0.028 per °C, so a Tris buffer made to pH 8.0 at 25 °C reads about 8.6 at 4 °C. The presets are 25 °C values.